Titrimetric Analysis of Aspirin
Name: Andrew Holmes
Student Number: 2103086
Course: BSc Forensic Science
Unit: Introduction to Forensic Chemistry
The purpose of the experiment was to determine the percentage of acetyl salicylic acid in an aspirin tablet and to compare this with the amount specified on the label. The average weight of one tablet before titration was found to be 328 mg. The amount of aspirin calculated from the results obtained from titrimetric analysis to be 317 mg. The stated dose of the tablets was 300 mg The percent purity of an aspirin tablet to the amount specified on the label was then calculated to be 96.47 % pure.
Aims and Objectives
The objectives of this experiment are to determine the percent aspirin in an aspirin tablet and to compare this with the amount specified on the label. This is carried out using an acid – base titration.
The aims of the experiment were as follows:
- Calculate the weight of one aspirin tablet.
- Calculate the molarity of diluted and stock NaOH solutions.
- Calculate the no. Of moles used in hydrolysis.
- To determine the percent aspirin in an aspirin tablet and compare it with the amount specified on the label.
- Calculate the percent purity of aspirin in one aspirin tablet to the specified amount on the label.
A titration is the basis of titrimetiric analysis. It is in which a solution containing a known concentration of a reagent is added to a solution containing an unknown concentration of a substance (In this case the sodium hydroxide remaining after hydrolysing acetyl salicylic acid) until the reaction between the reagent and the substance is complete. The solution of known concentration is called a standard solution and is the titrant (In this case hydrochloric acid). The volume of titrant added is measured and since its molarity is known the exact molarity of the reagent can be calculated
The titrant is added until the reaction is just complete at a point called the end point of the titration. An indicator is added to show when the end point has been reached; an indicator is some substance which changes colour at the end point and therefore show when the titration should be stopped.
The indicator being used in this experiment is phenolphthalein. Phenolphthalein is a weak acid, which is ionised in solution:
The acid itself is colourless but the anion P- is red. If the indicator is placed in an acidic solution where the hydrogen ion concentration is comparatively high, the dissociation of the indicator is suppressed. The unionised acid is colourless and therefore the solution is also colourless. IN alkaline solutions the hydroxyl ions remove the hydrogen ions produced in the above reaction, hence the concentration of P- ions is increased and the solution is given a pink colour.
This is a preview of the whole essay
The change from colourless to pink starts at pH = 8.2 and is complete by pH = 9.8. It is ideal for this experiment as its stoichiometeric point is near 9 which is ideal for a titration using a weak acid (acetyl salicylic acid) and a strong base (sodium hydroxide)
Aspirin is an analgesic drug that can be harmful if taken in excess. The main constituent is 2-ethnoylxybenzoic acid (acetyl salicylic acid.) The amount of aspirin in a sample can be quantified using its acidic properties in an acid-base reaction.
A known mass of aspirin is hydrolysed with a known excess amount of sodium hydroxide. The number of moles of NaOH left over is found by titrating with standard acid. The number of moles of NaOH required for the hydrolysis step can then be calculated because we know how many mole were originally present. Using the relationship given in the equation:
1 mole aspirin is hydrolysed by 2 moles NaOH
It is then possible to calculate the number of moles of acetyl salicylic acid that has been hydrolysed. This answer can then be compared to the label stating the aspirin content.
A sample of aspirin tablets, a stock solution of sodium hydroxide approximately 1 mol dm –3 and a standardised solution of hydrochloric acid of 0.1 mol dm –3 are supplied.
The determination of the exact molarity of the sodium hydroxide solution is carried out by diluting the supplied stock solution by a factor of ten. This is completed by adding 25 cm3 of the stock solution of sodium hydroxide to a volumetric flask and then making up to 250 cm3 with distilled water. 25 cm3 of the diluted stock solution is transferred into a conical flask and is titrated against the standard solution of 0.1 mol dm-3 hydrochloric acid. Phenolphthalein is used as an indicator. Before the titration can begin it is important that the solution in the burette (hydrochloric acid) is run though for a few seconds to ensure no air is contained in the burette, as this would compromise the results. The starting point of the hydrochloric acid in the burette is then recorded. The end point is reached when the pink colouring of the indicator is reduced so it is a clear liquid. The results of the titration are shown in table 1.
Approximately 1.3 to 1.7 g of aspirin tablet was accurately weighed out in a conical flask using an analytical balance. The conical flask was weighed before hand. Using a pipette 25 cm3 stock sodium hydroxide that is approximately 1.0 mol dm-3 was added to the conical flask and the aspirin tablets. A further 25 cm3 of distilled water was also added. The conical flask and its contents simmered gently over a Bunsen burner for 10 minutes. This was done to hydrolyse the acetyl salicylic acid.
The solution was then allowed to cool then transferred quantitivly into a 250 cm3 volumetric flask. The conical flask is rinsed out 2 or 3 times to ensure that all of the acetyl salicylic acid has been transferred to the volumetric flask. The solution in the volumetric flask is then made up to 250 cm3 using distilled water.
25 cm3 of the hydrolysed solution is pipetted into a clean conical flask. 15 drops of phenolphthalein was added to act as an indicator. The solution was then titrated with standard 0.1 mol dm-3 hydrochloric acid. The hydrochloric acid is then released from the burette slowly ensuring the hydrolysed solution is mixed constantly. The starting point of the hydrochloric acid in the burette must be recorded before the titration can begin. The end point is reached when the pink colouring of the indicator is reduced so it is a clear liquid. The titrations are repeated until the titration data is reproducible.
The apparatus must be cleaned and dried after each titration to ensure accurate results are achieved. Serval safty precautions must also be practiced. A lab coat must be worn at all times during the lab and long hair tied back. Also safety goggles must be worn when heating the acetyl salicylic acid and the sodium hydroxide.
1. Standardisation of 25cm3 sodium hydroxide solution.
Average volume of 0.1 mol dm-3 hydrochloric acid used = 24.98 2 d.p
2. Label information for the aspirin.
Stated dose of one tablet = 300 mg
Average weight of one tablet = 328 mg
Other ingredients; Maize starch
3. Weight of aspirin taken for analysis.
Note: Five tablets were used.
4. Titration of 25 cm3 hydrolysed solution.
Average volume of 0.1 mol dm-3 hydrochloric acid used = 7.35 2 d.p
In order to determine the percentage aspirin in one tablet and the mass of aspirin in one tablet to compare to the stated dose several calculations have to be completed. These are;
- Calculate molarity of the diluted stock NaOH solution.
- Calculate the number of moles of NaOH taken for hydrolysis.
- Calculate the number of moles of HCl used.
- Calculate the number of moles of NaOH used in hydrolysis.
- Calculate the number of moles of aspirin.
- Calculate the weight of aspirin.
- Calculate the weight of aspirin in one tablet.
- Calculate the percentage purity of the tablets.
Molecular Weight of Aspirin
1. Molarity of Diluted Stock NaOH Solution.
- Number of moles of NaOH taken for hydrolysis.
- Number of moles of HCl used.
- Number of moles of NaOH used in hydrolysis.
- Number of moles of aspirin.
- Weight of aspirin.
- Weight of aspirin in one tablet.
- Percentage purity of the tablets.
In this experiment a known mass of aspirin was hydrolysed with a known excess amount of sodium hydroxide. The number of moles of NaOH left over was found by titrating with a standard acid. The number of moles of NaOH required for the hydrolysis step was then calculated because we knew how many moles were originally present. Using this relationship given in the equation:
1 mole aspirin is hydrolysed by 2 moles NaOH
It was then possible to calculate the number of moles of acetyl salicylic acid that has been hydrolysed. The answer was then compared to the label stating the aspirin content.
The results obtained from titrimetric analysis of aspirin show that the percent purity of one tablet is 96.47% pure. It was also found that there was 317 mg of aspirin in the aspirin tablets compared to the 300 mg stated dose on the label. This is possibly due to experimental error, which is explained in further detail later. The average weight of one tablet was also found to be 328 mg the difference between the aspirin content and this value is due to the fact there is also maize starch added to the aspirin tablet to aid in digestion.
Experimental error in the results may have been due to the use of inaccurate glassware, human error, systematic error in the measurements that will influence the calculations made from the data obtained, glassware not properly rinsed and dried before use, dirt and air on the analytical balance, or the acetyl salicylic acid not fully hydrolysed.
If the experiment was to be repeated there are several changes that could be made to improve the accuracy and precision of the experiment. Accuracy and precision of the measurements taken in the experiment can influence the data obtained. The precision of a measurement refers to how close to one another these repeated measurements are. The accuracy of a series of measurements is the closeness of their average value to the true value.
If the experiment was to be repeated again it would be more accurate to use an indicator that changes colour around seven as this is where the equivalence point theoretically should be with a strong acid - strong base titration. Therefore in future the indictor which I believe would give more accurate and precise results would be bromothymol blue as this indicator changes colour at a pH of 6.0 and finishes at 7.6. The colour change is yellow to blue.
However this indicator would not produce more accurate results for the titration between the HCl and the hydrolysed solution as the NaOH is partly neutralised and thus is a more neutral solution. So the titration is more like a strong acid – weak base titration and so the theoretical end point is going to be lower therefore an indictor like methyl red which changes colour at 4.2 – 6.3 (from red to yellow) would produce the most accurate and precise results.
To summarise the points made in the discussion:
- The results obtained from titrimetric analysis of aspirin show that the percent purity of one tablet is 96.47% pure.
- The ratio of sodium hydroxide to hydrochloric acid is 1:1; therefore, the molarity of diluted and stock sodium hydroxide solutions can be calculated by working out the number of moles of hydrochloric acid used.
- The number of moles of sodium hydroxide used in hydrolysis can be calculated by subtracting the number of moles of sodium hydroxide taken for hydrolysis by the number of moles of sodium hydroxide remaining after hydrolysis.
- The weight of one aspirin tablet was calculated by working out the molecular weight of aspirin and multiplying it by the number of moles of aspirin used, to find the weight of the sample of aspirin used in the experiment; and then dividing it by the number of tablets in the sample.
- In an acid-base titration, as a strong acid is gradually added to a strong base, it reacts with the base and the pH gradually falls.
- Accuracy and precision of the measurements taken in the experiment can influence the data obtained. Experimental error in the results may have been due to the use of inaccurate glassware, human error, systematic error in the
- Measurements that will influence the calculations made from the data obtained, glassware not properly rinsed and dried before use, dirt and air on the analytical balance, or the acetyl salicylic acid not fully hydrolysed.
- The correct safety precautions were adhered to during experimentation. These included handling glassware in a careful and appropriate manner, wearing safety goggles during the heating stage of the experiment, and using tongs and a heatproof mat.
Jones, L; Atkins, P. (1999) Chemistry: Molecules, Matter and Change 4th edition. America: Freeman